EXCHANGE 


The  Electromotive  Force  and  Free 

Energy  of  Dilution  of  Lithium 

Chloride  in  Aqueous  and 

Alcoholic  Solutions 


A  DISSERTATION 


SUBMITTED  TO  THE  FACULTY  OF  THE  GRADUATE  COLLEGE  OF  THE 

STATE  UNIVERSITY  OF  IOWA  IN  PARTIAL  FULFILLMENT  OF 

THE  REQUIREMENTS  FOR  THE  DEGREE  OF 

DOCTOR  OF  PHILOSOPHY 


WTHt 


By 

FRANKLIN  SPENCER  MORTIMER 
1917 


BASTON,  PA.: 
BSCHSNBACH    PRINTING    COMPANY 


The  Electromotive  Force  and  Free 

Energy  of  Dilution  of  Lithium 

Chloride  in  Aqueous  and 

Alcoholic  Solutions 


A  DISSERTATION 


SUBMITTED  TO  THE  FACULTY  OF  THE  GRADUATE  COLLEGE  OF  THE 

STATE  UNIVERSITY  OF  IOWA  IN  PARTIAL  FULFILLMENT  OF 

THE  REQUIREMENTS  FOR  THE  DEGREE  OF 

DOCTOR  OF  PHILOSOPHY 


By 


FRANKLIN  SPENCER  MORTIMER 

i\ 

1917 


EASTON,  PA.: 

ESCHENBACH    PRINTING    COMPANY 
1918 


ACKNOWLEDGMENTS. 

The  author  wishes  to  express  his  sincere  appreciation  for  the  assist- 
ance and  inspiration  of  Dr.  J.  N.  Pearce,  at  whose  suggestion  and  under 
whose  direction  this  investigation  was  made. 

Thanks  are  also  extended  to  Dr.  E.  W.  Rockwood,  whose  hearty  co- 
operation in  obtaining  apparatus,  has  aided  materially  in  this  work. 

The  writer  also  desires  to  thank  Dr.  W.  J.  Karslake  and  Prof.  A.  W. 
Hixon  for  their  valuable  instruction  and  kindly  interest. 

F.  S.  M. 


The   Electromotive   Force  and   Free   Energy 

of    Dilution   of   Lithium   Chloride   in 

Aqueous  and  Alcoholic  Solutions. 


The  study  of  the  electromotive  forces  in  aqueous  solutions  of  electro- 
lytes has  always  presented  an  attractive  field  for  research.  One  has  but 
to  scan  the  literature1  of  the  last  twenty-five  years  to  be  convinced  of  the 
enormous  amount  of  work  which  has  been  done  in  this  field  alone.  It  is 
only  in  more  recent  years  that  attempts  have  been  made  to  extend  electro- 
metric  measurements  with  the  view  of  correlating  the  results  of  the  measure- 
ments of  electromotive  force  of  concentration  cells  with  other  colligative 
properties  of  solutions. 

Dolezalek2  has  determined  the  electromotive  forces  at  30°  of  hydrogen- 
chlorine  cells,  using  hydrochloric  acid  for  the  electrolyte  and  in  concentra- 
tions ranging  from  4.98  N  to  12.25  N.  His  results  are  sufficiently  accurate 
to  permit  the  calculation  of  the  free  energy  of  transference  of  hydrochloric 
acid  for  the  various  concentrations  used. 

Using  concentration  cells  of  the  type 

Ag  |  AgCl,  HCki  |  HCfc,,  AgCl  |  Ag 

Jahn3  determined  with  great  accuracy  the  electromotive  forces  for  concen- 
trations ranging  from  0.033  N  to  0.00167  N.  Similar  cells  were  made  using 
correspondingly  dilute  solutions  of  the  chlorides  of  sodium  and  potassium. 
From  the  results  obtained  he  concluded:  (i)  that  the  mobility  of  the  ions 
is  dependent  to  a  high  degree  upon  their  concentration,  increasing  with 
increasing  concentration,  (2)  that  the  conductivity  of  an  electrolyte  is  not 
a  true  measure  of  its  dissociation. 

Tolman  and  Ferguson4  determined  the  electromotive  forces  for  concen- 
tration cells  of  the  type  Hg  |  HgCl,  HC1 1  HC1,  H2 1  Pt  at  18°  and  combined 
so  as  to  eliminate  transference.  From  these  they  calculated  the  free  energy 
of  dilution  of  hydrochloric  acid,  and  the  ratios  of  the  fugacities  of  the  ions 
and  of  the  molecules.  They  found  that  for  strong  electrolytes,  even  in 
dilute  solutions,  the  fugacity,  or  activity,  of  the  ions  is  not  strictly  propor- 
tional to  the  concentration,  while  that  of  the  undissociated  electrolyte  is 
very  far  from  proportional  to  its  concentration. 

1  For  a  complete  summary  see  Abegg-Auerbach  and  Luther,  "Messungen  elektro- 
motorischer  Krafte." 

2  Z.  physik.  Chem.,  26,  321  (1898). 

3  Ibid.,  33,  545  (1900). 

4  J.  Am.  Chem.  Soc.,  34,  232  (1912). 


Mclnnes  and  Parker1  have  extended  the  electrometric  measurement  of 
the  free  energy  of  dilution  to  aqueous  solutions  of  potassium  chloride. 
Using  silver  chloride  and  amalgam  electrodes,  they  determined  the  elec- 
tromotive force  of  cells  both  with  and  without  transference.  From  their 
data  they  were  able  to  calculate  the  transport  numbers  and  the  activity 
ratios  of  the  ions.  They  found  that  the  concentration  ratios  calculated 
from  conductivity  are  invariably  higher  than  the  activity  ratios  determined 
by  the  electromotive-force  method;  further,  that  the  activity  ratio  ap- 
proaches the  value  of  the  concentration  ratio  as  the  dilution  increases. 

Ferguson2  also  has  worked  with  concentration  cells  containing  hydro- 
chloric acid,  using  electrodes  reversible  to  both  ions  in  cells  with  and  with- 
out transference.  He  also  found  that  the  activity  ratios  are  less  than  the 
concentration  ratios. 

Somewhat  similar  results  were  obtained  by  Ellis3  for  hydrochloric  acid 
solutions  in  concentrations  ranging  from  0.00167  N  to  4.484  N.  Assuming 
provisionally  that  the  activity  coefficient  of  the  most  dilute  solution  is 
substantially  equal  to  the  ionization  coefficient  derived  from  the  conduc- 
tance ratio,  he  calculated  a  series  of  absolute  activity  coefficients  for  the 
remaining  concentrations.  Beginning  with  the  most  dilute  solution,  the 
activity  coefficients,  thus  calculated,  decrease  with  increasing  concen- 
tration, pass  through  a  minimum  at  0.5  N  and  then  increase  with  further 
increase  in  concentration.  The  conductance- viscosity  ratio  Xrj/Xo^o  de- 
creases throughout  with  increase  in  the  concentration  of  the  acid. 

Thus  far  water  has  been  the  only  solvent  used  in  investigations  of  this 
nature.  It  was  thought  worth  while  therefore,  to  make  a  systematic  com- 
parative study  of  the  influence  of  the  solvent  upon  the  electromotive  forces, 
the  transport  numbers,  the  activity  ratios  of  the  ions  and  molecules,  and 
the  free  energy  of  dilution  of  electrolytes  in  water  and  organic  solvents. 

The  concentration  cells  chosen  for  this  work  consist  of  solutions  of  lithium 
chloride  in  water  and  in  the  five  lower  alcohols  of  the  paraffin  series. 

In  such  a  study  it  is  necessary  to  make  use  of  two  types  of  concentration 
cells.  Let  us  consider  first  the  cell  involving  transference,  e.  g., 


Ag 


AgCl,  LiCl 


LiCl,  AgCl 


Ag, 


{' 

During  the  passage  of  one  faraday  of  electricity,  one  equivalent  of  the 
chloride  ion  is  formed  on  the  dilute  side  from  the  silver  chloride  electrode, 
while  on  the  concentrated  side  one  equivalent  of  the  chloride  ion  is  removed 
from  the  solution  to  the  electrode.  At  the  same  time  Nc  equivalents  of 
the  lithium  ion  migrate  into  the  dilute  chamber  and  (i  —  Nc)  equivalents 
of  the  chloride  ion  migrate  to  the  concentrated  side.  The  total  result  is 

1  J.  Am.  Chem.  Soc.,  37,  1445  (1915). 

2  J.  Phys.  Chem.,  20,  326  (1916). 

3  J.  Am.  Chem.  Soc.,    38,  737  (1916). 


the  transfer  of  Nc  equivalents  of  lithium  chloride  from  the  concentrated 
to  the  dilute  solution,  or  from  the  solution  of  activity  £"  to  that  of  £'. 
The  free  energy  accompanying  the  transfer  of  one  mole  of  the  salt  is  given 
by  the  relation 


where  E  is  the  electromotive  force,  F  the  faraday  (96494  coulombs),  Nc 
the  transport  number  of  the  cation,  R  the  gas  constant  (8.316  j.),  T  the 
absolute  temperature  (298.09°). 

The  following  cell  does  not  involve  transference  : 


Ag 


AgCl,  LiCl 


LiCl  (Li.Hg,)  —  (Li.HgJ  LiCl 

r  r 


LiCl,  AgCl 

r 


Ag 


The  passage  of  one  faraday  of  electricity  involves,  on  the  dilute  side  only, 
the  formation  of  one  equivalent  of  lithium  chloride  from  the  silver  chloride 
and  the  amalgam  electrodes.  On  the  concentrated  side  there  is  transferred 
to  the  electrodes  from  the  solution  one  equivalent  of  lithium  chloride.  The 
free  energy  accompanying  this  change  is 

ge.  (2) 


Combining  Equations  i  and  2  we  arrive  at  an  expression  for  calculat- 
ing the  transport  number  of  the  cation  directly  from  electromo.tive-force 
measurements,  viz., 

Nc  =  f  ,  (3) 

•&1 

All  cells  on  closed  circuit  tend  to  operate  until  the  activities  of  the  two 
solutions  become  equal.  In  cells  without  transference  such  an  equaliza- 
tion by  direct  diffusion  of  the  molecules  and  ions  is  impossible.  The  same 
result  is  obtained  by  the  formation  of  the  salt  from  the  electrodes  on  the 
dilute  side  and  the  simultaneous  removal  of  the  salt  to  the  electrodes  on 
the  concentrated  side.  It  is  obvious,  therefore,  that  the  free  energy  of 
dilution  of  lithium  chloride  is  equal  to  the  sums  of  the  free  energy  of  dilu- 
tion of  the  separate  ions,  i.  e., 


-AF,  =  E,F  =  RT  log,  =  RT  log. 

£  (LiCl)  £  Li*-?  Cl~ 

Assuming  that  £"Li+  =  £"C1-  and  that  £'Li+  =  £'Ci-,  then  for  the  chloride  ion 


=  RT  log,          l>  =  2RT  log,     ar. 


If  this  assumption  is  not  true,  and  most  probably  it  is  not,  then  the  calcu- 
lated ratio  of  the  activities  of  the  chlorine  ions  may  be  taken  as  the  ratio 
of  the  square  roots  of  the  products  of  the  activities  of  the  two  ions. 

The  free  energy  of  dilution  in  joules  per  gram-equivalent  weight  is  equal 


8 

to  EI,  the  electromotive  force  without  transference,  multiplied  by  the  value 
of  the  faraday,  or  96494  coulombs. 

Materials  and  Apparatus. 

The  conductivity  water  used  throughout  the  work  was  prepared  accord- 
ing to  the*method  of  Jones  and  Mackay.1 

Ethyl  Alcohol. — Ordinary  95%  alcohol  was  allowed  to  stand  over  fresh 
quicklime  for  three  weeks;  it  was  then  decanted  and  distilled.  The  distil- 
late was  allowed  to  stand  over  anhydrous  copper  sulf ate  for  one  week  and 
then  redistilled.  This  distillate  was  refluxed  over  metallic  calcium  for 
ten  hours  and  again  distilled.  To  the  last  distillate  were  added  a  few  crys- 
tals of  pure  dry  silver  nitrate  and  it  was  then  refluxed  for  two  hours  to  re- 
move reducing  agents.  The  distillate  from  this  treatment  was  collected 
and  preserved  in  dry  glass-stoppered  bottles,  being  protected  from  the  air 
during  distillation  by  a  drying  train  containing  fused  calcium  chloride. 
In  each  distillation  a  fractionating  column  was  used  and  only  that  middle 
portion  passing  over  between  77.9°  and  78°  (uncorr.)  was  retained. 

The  remaining  alcohols,  viz.,  methyl,  w-propyl,  w-butyl  and  isoamyl, 
were  of  Kahlbaum's  best  grade.  These  were  further  purified  in  the  same 
manner  as  was  the  ethyl  alcohol,  except  that  the  treatment  with  quicklime 
was  omitted.  The  uncorrected  distillation  temperatures  were:  methyl, 
64.9-65.1°;  n-propyl,  95-7-95-9°;  w-butyl,  115.8-116.2°;  isoamyl, 
129.9-136.2°. 

Lithium  Chloride. — Kahlbaum's  best  grade  of  lithium  chloride  was  re- 
crystallized  four  times  by  passing  pure  hydrogen  chloride  into  a  hot 
saturated  solution  of  the  salt  in  conductivity  water.  The  crystals  were 
filtered  on  a  Buchner  funnel  and  sucked  dry.  They  were  then  heated  in  a 
platinum  dish  in  an  electric  oven  in  which  the  temperature  was  gradually 
raised  to  150°.  The  dry  salt  was  then  finely  powdered  in  a  hot  agate 
mortar  and  transferred  to  porcelain  boats.  These  were  placed  in  a  com- 
bustion tube  and  heated  for  several  hours  at  1 60°  in  a  rapid  stream  of  dry 
hydrogen  chloride.  All  traces  of  the  latter  were  removed  by  a  stream  of 
dry  hydrogen,  after  which  the  boats  were  allowed  to  slide  quickly  into  large 
glass-stoppered  weighing  tubes. 

Solutions. — All  of  the  solutions  were  prepared  by  first  dissolving  an 
amount  of  the  salt  in  excess  of  that  required  for  the  highest  concentration. 
The  chloride  content  was  then  determined  in  at  least  four  separate  samples 
by  the  Drechsel2  modification  of  the  Volhard  method.  All  of  the  various 
concentrations  in  any  given  solvent  were  made  by  the  proper  dilution  of 
this  mother  solution,  the  greatest  care  being  taken  to  secure  the  least  pos- 
sible contact  with  the  air.  All  of  the  measuring  apparatus  was  certified 
and  the  solutions  were  made  up  to  volume  at  25°. 

1  Am.  Chem.  J.,  19,  83  (1897). 

2  Z.  anal.  Chem.,  16,  351  (1877). 


Mercury. — The  mercury  used  was  repeatedly  washed  by  spraying  through 
dilute  nitric  acid,  and  then  further  purified  by  distilling  in  a  current  of  air 
under  reduced  pressure. 

Lithium  Amalgam. — This  was  prepared  by  the  electrolysis  of  a  saturated 
solution  of  the  salt  in  pyridine,  using  pure  redistilled  mercury  as  the 
cathode.  It  was  then  washed  with  absolute  alcohol,  quickly  dried  by 
gentle  heating  under  reduced  pressure  (Gaede  pump),  then  filtered  through 
a  fine  capillary  tube  into  a  sealed  glass  container  from  which  the  air  had 
previously  been  displaced  by  dry  hydrogen. 

Electrodes. — The  silver  chloride  electrodes  consisted  of  short,  thick 
pieces  of  pure  silver  wire  fused  into  the  ends  of  glass  tubes.  To  the  ends 
within  the  tubes  were  soldered  long  copper  wires  which  were  of  such  length 
that  they  could  be  bent  into  small  mercury  cups,  thus  making  contact  with 
the  wire  leads.  Twelve  or  fifteen  of  the  electrodes  thus  prepared  were 
first  grouped  as  cathodes  about  a  single  pure  silver  anode  immersed  in  a 
solution  of  potassium-silver-cyanide.  After  a  dense  white  coating  of  silver 
had  been  formed  they  were  removed,  rinsed  and  then  inserted  as  anodes  in 
a  i  .o  AT  hydrochloric  acid.  On  passing  a  current  from  a  single  accumulator 
for  one  to  two  minutes  there  is  formed  a  closely  adhering  reddish  brown 
deposit  of  silver  chloride.  For  any  one  series  of  measurements  the  chloride 
electrodes  were  always  first  checked  against  each  other.  This  was  done 
by  grouping  them  in  a  dilute  solution  of  hydrochloric  acid  and  observing 
the  potential  difference  between  each  electrode  and  another  taken  as  stand- 
ard. Only  those  varying  by  less  than  0.05  mv.  were  chosen.  Three  of 
these  electrodes  were  placed  in  each  half  cell.  After  they  had  been  in 
contact  with  each  other  for  four  or  five  hours  the  potential  readings  were 
taken.  In  taking  these  readings  the  electrodes  in  each  half -cell  were  first 
checked  against  each  other.  Unless  at  least  two  electrodes  differed  by  less 
than  0.02  mv.,  the  cell  was  disconnected  and  the  electrodes  replated. 
When  the  potential  deviations  proved  to  be  within  these  limits  the  poten- 
tials were  read  between  each  single  electrode  of  one  cell  and  the  separate 
electrodes  of  the  other  cell  and  the  mean  value  taken.  For  a  further 
check  the  electrodes  in  each  half  cell  were  connected  and  the  electromotive 
force  between  the  combined  electrodes  was  then  determined. 

The  form  of  cell  adopted  was  such  that  concentration  cells,  both  with  and 
without  transference,  could  be  obtained  from  a  single  set-up  of  the  ap- 
paratus. Two  half -cells  of  the  form  used  by  Ferguson,1  each  having  two 
side  tubes  and  containing  solutions  of  the  desired  concentrations  in  con- 
tact with  the  chloride  electrodes,  were  so  arranged  that  from  one  set  of 
side  tubes  a  liquid  junction  could  be  made  and  the  electromotive  force 
with  transference  thus  obtained.  The  other  set  of  tubes  were  thus  left 

1  Loc.  cit. 


free  for  liquid  connection  with  small  cells  into  which  dipped  the  amalgam 
electrodes,  thus  forming  the  cell  without  transference. 

Liquid  connection  between  the  half  cells  was  effected  by  means  of  an 
inverted  T-tube  fitted  with  a  three-way  stopcock.  To  minimize  diffusion 
loose  plugs  of  cotton  were  inserted  into  the  bore  of  these  stopcocks.  For 
cells  with  transference  it  is  very  essential  that  a  sharp  boundary  be  pro- 
duced between  the  solutions  immediately  before  measurements  are  taken. 
Fresh  contacts  were  readily  made  by  drawing  more  of  each  solution  into 
the  free  limb  of  the  T-tube.  It  was  found  that  the  electromotive  forces 
of  cells  directly  connected  remained  constant  for  several  days  when  this 
precaution  was  observed. 

The  amalgam  electrodes  were  similar  to  those  used  by  Maclnnes  and 
Parker;1  the  method  of  procedure  was  also  the  same.  As  a  result  of 
numerous  experiments,  a  concentration  of  0.002  %  was  found  to  give  the 
best  results.  It  was  found  that  by  dropping  from  20  to  30  drops  per  minute 
no  appreciable  bubbling  occurred  and  the  voltage  remained  very  constant 
for  several  minutes.  Occasionally,  the  galvanometer  would  waver  slightly 
but  it  would  immediately  return  to  zero  upon  the  formation  of  another 
drop  of  the  amalgam. 

All  measurements  of  the  electromotive  force  were  made  with  a  Wolff 
potentiometer  in  connection  with  a  sensitive  Leeds  and  Northrup,  "Type 
H,"  wall  galvanometer.  A  Cadmium- Weston  cell  which  had  been  recently 
standardized  and  frequently  rechecked  against  a  similar  element,  also 
recently  standardized  by  the  Bureau  of  Standards,  was  used  as  the  standard 
of  reference.  Although  its  temperature  coefficient  is  practically  negligible, 
this  cell  was  kept  in  an  insulated  glass  beaker  suspended  in  the  constant 
temperature  bath.  All  measurements  were  made  at  25°  ±0.01°. 
Precision  and  Duplication  of  Results. 

The  resistance  of  the  solutions  increases  with  increasing  molecular  weight 
of  the  solvent ;  obviously,  the  precision  to  be  obtained  decreases  somewhat 
accordingly.  However,  it  was  not  found  difficult  to  duplicate  potentiom- 
eter readings  for  any  pair  of  concentrations  in  any  of  the  solvents  used. 

Quadruplicate  determinations  of  the  electromotive  force  were  made  for 
all  pairs  of  concentrations  in  water,  methyl  alcohol  and  ethyl  alcohol. 
This  involved  four  distinctly  different  set-ups  of  the  apparatus,  including 
new  solutions  and  three  freshly  plated  concordant  electrodes  for  each  half 
cell.  The  results  recorded  represent,  therefore,  the  mean  of  four  "mean" 
potential  differences  which,  over  all,  do  not  differ  by  more  than  0.05  mv. 

The  electromotive  forces  recorded  for  solutions  in  w-propyl  and  isoamyl 

alcohols  are  the  mean  of  duplicate  measurements,  made  as  previously 

described.     This  was  deemed  sufficient,  since  the  potentiometer  readings 

checked  exactly  to  the  last  significant  figure.     The  supply  of  w-butyl 

1  Loc.  tit. 


II 

alcohol  at  our  disposal  permitted  but  a  single  determination  of  the  electro- 
motive force  for  a  single  pair  of  concentrations  in  this  solvent.  Since, 
however,  each  half  cell  contained  three  electrodes  which  checked  over  all 
to  0.02  mv.,  and  since  each  electrode  was  measured  directly  against  each 
of  the  three  electrodes  in  the  opposite  cell,  the  results  obtained  for  this 
solvent  may  be  considered  as  the  mean  of  six  different  determinations  for 
the  same  pair  of  solutions. 

Discussion. 

Lithium  chloride  was  chosen  as  the  electrolyte  because  of  its  relatively 
high  solubility  in  all  of  the  solvents  under  consideration.  Its  use  is  ham- 
pered, however,  by  certain  disadvantages.  Various  lines  of  evidence  point 
to  its  inherent  tendency  to  polymerize  when  dissolved  in  the  alcohols  and 
other  organic  solvents,  l  and  also  to  the  tendency  of  both  the  molecules  and 
the  ions  to  form  complex  solvates  with  the  solvents.2  Washburn  and  Mac- 
Innes3  have  calculated  for  a  0.5  N  aqueous  solution  of  lithium  chloride  a 
probable  hydration  of  18  moles  of  water  to  one  mole  of  the  salt. 

For  a  given  temperature,  the  complexity  of  these  solvates  must  increase 
with  increase  in  the  proportion  of  the  solvent.  When  the  phenomenon 
of  association  is  present,  it  is  very  probable  also  that  both  the  polymerized 
molecules  and  their  ions  are  solvated  to  some  extent.  For  any  given  con- 
centration of  the  salt  in  any  one  of  the  solvents,  except  perhaps  water,  we 
must  take  into  consideration  the  possible  existence  of  an  equilibrium  not 
only  between  the  solvated  simple  and  complex  molecules,  but  also  between 
the  two  molecular  species  and  their  solvated  ions,  e.  g., 

l4+.Si+LiCl-2.S5  (a) 


-  52^±UCl.S,^±(LiCl),.S4  or  (5) 

^LijCl+.S6+Cl-.Si  (6) 
where  Si,  S2,  Ss,  etc.,  represent  the  number  of  moles  of  combined  solvent. 

Slight  association,  if  any,  should  be  expected  in  water,  methyl  alcohol 
and  ethyl  alcohol.  For  the  alcohols,  the  relative  degree  of  association  of 
the  salt  should  increase  with  the  decrease  in  the  dissociating  power  of  the 
solvent,  or,  as  in  the  present  work,  with  an  increase  in  the  molecular  weight 
of  the  alcohols. 

An  increase  in  the  concentration  of  the  salt  will  tend  to  displace  the 
above  equilibrium  toward  the  right.  With  an  increase  in  the  concentra- 
tion of  the  polymerized  salt  molecules  there  is  very  probably  associated 
also  an  increase  in  the  formation  of  complex  ions  of  greater  electro-affinity.4 

1  Andrews  and  Ende,  Z.  phys.  Chem.,  17,  136  (1905). 

2  Jones  and  Getman,  Ibid.,  46,  261  (1903);  Am.  Chem.  /.,  32,  338  (1904);  Jones  and 
McMaster,  Ibid.,  35,  445  (1906). 

3  /.  Am.  Chem.  Soc.,  33,  1705  (1911). 

4  Sachanov,  /.  Russ.  Phys.  Chem.  Soc.,  44,  324  (1912);  Pearce,  /.  Phys.  Chem.,  19, 
14  (1915)- 


12 

On  the  other  hand,  a  sufficient  decrease  in  the  concentration  will  naturally 
result  in  the  formation  of  the  solvated  simple  molecules  and  their  ions. 
While  the  magnitude  of  the  solvation  per  mole  of  the  solute  decreases  with 
increase  in  concentration,  the  total  amount  of  solvent  removed  from  action 
as  combined  solvent  will  be  greater,  the  greater  the  concentration  of  the 
salt.  With  a  decrease  in  the  amount  of  effective  solvent,  due  to  solvation, 
there  is  an  abnormal  increase  in  the  actual  concentration  of  the  solute 
particles,  and  very  probably,  therefore,  a  corresponding  abnormal  increase 
in  the  activity  of  both  the  molecules  and  the  ions,  whether  simple  or  com- 
plex. It  is  obvious  therefore  that  for  such  equilibria  as  that  represented 
above  we  should  expect  anomalous  relations  to  appear  as  we  pass  from 
very  dilute  to  very  concentrated  solutions. 

Changes  in  the  extent  of  polymerization,  in  ion  complexity  and  in  solva- 
tion with  change  in  concentration  must  abnormally  effect  the  electromotive 
forces  of  cells  without  transference.  These  changes  together  with  the 
accompanying  changes  in  viscosity  must  also  affect  the  mobilities  of  the 
ions  and,  therefore,  the  electromotive  forces  of  cells  with  transference. 

Solvation  and  polymerization  of  the  salt  are  not  the  only  factors  involved 
in  the  study  of  solutions  of  electrolytes  in  organic  solvents.  The  activity 
of  the  dissolved  molecules  and  ions  is  a  function  of  their  concentration. 
According  to  the  Nernst-Thomson  rule,  the  dissociating  power  of  the  sol- 
vent will  be  greater,  the  greater  is  its  dielectric  constant.  Consequently, 
considering  a  given  concentration  of  the  electrolyte  in  various  solvents, 
that  solution  will  contain  the  greatest  concentration  of  ions,  and  hence  the 
greatest  ion  activity,  whose  solvent  possesses  the  highest  dielectric  con- 
stant. 

Euler1  has  found  that  the  dielectric  constants  of  solutions  increase  with 
their  ion  concentration.  In  some  work  carried  out  in  this  laboratory 
several  years  ago2  it  was  found  that  the  dielectric  constant  of  solutions 
of  silver  nitrate  in  methyl  and  ethyl  alcohol  increases  linearly  with  the 
concentration  up  to  0.05  N. 

Walden3  has  found  that  the  dielectric  constants  of  solvents  of  feeble 
ionizing  power  are  increased  by  dissolving  in  them  certain  binary  salt. 
The  increase  in  the  value  of  the  constant  is  dependent  upon  the  constitution 
of  the  salt.  Strong  salts,  e.  g.,  lithium  chloride,  when  dissolved  possess 
high  dielectric  constants  and  exhibit  a  great  tendency  to  ionize.  The 
degree  of  ionization  of  the  salt  depends  both  on  the  dissociating  power  of 
the  solvent  and  the  ionizing  tendency  of  the  salt.  As  both  of  these  factors 
increase  with  the  dielectric  constant  the  highest  degree  of  ionization — 
hence  of  ion  concentration  and  ion  activity  for  a  given  concentration  of 

1  Z.  phys.  Chem.,  28,  619  (1899). 

2  Unpublished. 

3  J.  Am.  Chem.  Soc.,  35,  649  (1913)- 


13 

the  salt, — will  be  found  in  a  system  where  both  the  solvent  and  the  solute 
possess  large  dielectric  constants. 

TABLE  I. — ELECTROMOTIVE  FORCES  WITH  TRANSFERENCE  ( VOLTS). 

Concentration  ratios  (10  :  1). 


Solvent. 
Water  

1.0-0.1. 
0.0^194 

0.5-0.05. 
O.O^SO^ 

0.1-0.01. 

o  o^sS'? 

0.05-0.005. 

o  0^64.0 

Methyl  ale  
Ethyl  ale 

0.03857 
O    O^^2^ 

o  .  04004 
o  03560 

0.04104 
o  03820 

w-Propyl  ale  
w-Butyl  ale  

0.02740 

0.02817 

o  0271 

0.02845 
o  0242 

o  .  03  i  i 
o  0250 

Isoamyl  ale  

0.02630 

0.0255 

o  .  0243 

0.0240 

For  solutions  in  water,  methyl,  ethyl,  and  propyl  alcohols  the  electro- 
motive forces  with  transference  increase  regularly  with  increase  in  dilution. 
Normally  this  is  just  what  we  should  expect.  In  the  w-butyl  alcohol  the 
electromotive  force  first  decreases,  passes  through  a  minimum  and  then 
increases  slightly  upon  further  dilution.  For  the  isoamyl  alcohol  there 
is  a  corresponding  decrease  in  the  values  of  the  electromotive  force,  tending 
toward  a  minimum  in  the  most  dilute  pair. 

TABLE  II. — ELECTROMOTIVE  FORCES  WITHOUT  TRANSFERENCE  (VOLTS). 

Concentration  ratios  (10  :  1). 


Solvent. 

Water  

1.0-0.1. 

o.  1  14.  ^3 

0.5-0.05 
o  10868 

0.1-0.01. 
O    IO4.T* 

0.05-0.005. 
O   OOQS? 

Methyl  ale  
Ethyl  ale 

0.09387 

o  08875 

0.07978 

o  07170 

0.07162 
O  06  I  43 

w-Propyl  ale.  .  .  . 
w-Butyl  ale.     .  . 

o  0825 

0.07885 
O   O772 

0.05890 
o  0^70 

0.0575 

o  0501 

Isoamvl  ale.  . 

O.O7OI 

o  .  062  =; 

0.0=;^ 

O.O4.^O 

For  cells  without  transference  the  electromotive  force  of  all  pairs  in  the 
same  solvent  decreases  with  increase  in  dilution.  This  is  exactly  the  op- 
posite of  what  we  should  expect;  for,  with  a  normality  ratio  of  10  :  i,  it 
would  be  expected  that  the  concentration  ratios,  as  well  as  the  activity 
ratios,  of  the  ions  should  gradually  increase  to  the  value  of  10  at  infinite 
dilution.  It  is  evident  that  lithium  chloride  behaves  abnormally  in  all 
of  these  solvents.  A  discussion  of  this  abnormality  will  be  taken  up  later. 

Using  the  conductivity  data  of  Greene1  for  solutions  of  lithium  chloride 
in  water  and  those  of  Jones  and  Turner2  for  solutions  in  ethyl  alcohol, 
I  calculated  the  electromotive  forces  for  cells  both  with  and  without 
transference.  In  both  solvents  the  calculated  values  increase  with  in- 
creasing dilution,  but  in  neither  type  of  cell  does  the  Nernst  equation  apply 
even  approximately. 

The  transference  number  of  the  lithium  ion  increases  with  increase  in 
dilution  in  each  of  the  solvents  studied.  For  a  given  change  in  the  dilution 

1  Trans.  Chem.  Soc.,  93,  2042  (1908). 

2  Am.  Chem.  J.,  40,  558  (1908). 


this  increase  is  least  in  the  aqueous  and  greatest  in  the  ethyl  alcohol. 
Except  for  the  most  dilute  cell  in  the  ethyl  alcohol,  the  highest  values  for 
the  transference  number  are  obtained  for  solutions  in  methyl  alcohol. 
For  similar  concentrations  in  the  alcohols  the  transference  number  de- 
creases with  increase  in  the  molecular  weight  of  the  solvent  until  isoamyl 
alcohol  is  reached,  where  the  corresponding  values  are  found  to  increase. 

TABLE  III. — TRANSFERENCE  NUMBERS  OF  THE  LITHIUM  ION. 

Concentration  ratios  (10  :  1). 


Solvent. 

Water 

1.0-0.1. 
O    27Q 

0.5-0.05. 
O    ^22 

0.1-0.01. 
O    ^4^ 

0.05-0.005. 

o  ^6s 

Methyl  ale 

O   4.1  1 

o  502 

O    ^7^ 

Ethyl  ale  

O    ^74 

O   4Q7 

o  622 

w-Propyl  ale 

O    ^57 

O   48^ 

O    ">4I 

w-Butyl  ale  

0.332 

O.35I 

o  418 

O.499 

Isoamyl  ale  

0.375 

0.408 

0-454 

0-555 

When  the  concentration  is  such  that  the  possibility  of  the  presence  of 
complex  ions  exists,  the  transference  number  obtained  for  an  ion,  e.  g., 
the  anion,  will  be  the  sum  of  the  transference  numbers  of  the  simple  and 
complex  anions  present.  Only  at  sufficiently  high  dilutions  will  we  obtain 
the  transference  number  of  the  simple  ion. 

The  values  for  the  transference  numbers  obtained  by  the  electromotive- 
force  method  are  the  mean  values  between  the  two  concentrations  con- 
stituting the  cell.  It  is  therefore  difficult  to  make  a  comparison  with  data 
directly  determined.  Table  IV  gives  the  values  determined  by  Kohl- 
rausch  and  Holborn1  for  the  transference  number  of  the  lithium  ion  in 
aqueous  solutions  at  25°.  The  mean  values  have  been  calculated  and  are 
inserted  for  comparison. 

TABLE  IV. — COMPARISON  OF  THE  TRANSFERENCE  NUMBERS  OF  THE  LITHIUM  ION. 

1.0-0.1  0.5-0.05.  0.1-0.01.  0.05-0.005. 

Kohlrausch 0.261     0.3100.270     0.3300.310     0.3700.330     O.39O2 

Mean  Nc  (K) 0.285  0.300  0.340  0.360 

Mean  Nc  (e.  m.  f.). . .  .          0.279  0.322  0.343  0.365 

The  agreement  is  as  satisfactory  as  could  be  expected  and  it   confirms  the 
applicability  of  this  method  of  determining  transference  numbers. 

TABLE  V. — FREE  ENERGY  OF  DILUTION  (JOULES). 

Concentration  ratios  (10  :  1). 


Solvent. 

Water  

1.0-0.1. 
1  1  03  2 

0.5-0.05. 

10487 

0.1-0.01. 
10067 

0.05-0.005. 
9608 

Methyl  ale.  .  . 

9205 

7698 

6911 

Ethyl  ale 

8564 

6919 

5928 

w-Propyl  ale 

7608 

5684 

CC4.8 

w-Butyl  ale.  .  . 

796l 

7450 

5587 

4834 

Isoamyl  ale.    . 

6764 

6031 

5162 

4149 

1  "Leitvermogen 
2  Extrapolated. 

der  Elektrolyte,"  p.  201. 

It  will  be  observed  that  for  a  given  normality  ratio  (10  :  i)  the  free 
energy  of  dilution  of  lithium  chloride  decreases  regularly  as  we  pass  from 
the  lower  to  the  higher  members  of  the  same  homologous  series.  Further- 
more, for  any  one  of  the  solvents  studied  the  free  energy  decreases  as  the 
pair  of  concentrations  involved  become  more  dilute.  Since  the  free  energy 
of  dilution  is  proportional  to  the  logarithm  of  the  ratio  of  the  activities  of 
the  chlorine  ions,  we  should  expect,  for  the  normality  ratio  used,  that  the 
free  energy  would  increase  to  a  maximum  at  infinite  dilution. 

Table  VI  gives  a  summary  of  the  activity  ratios  of  the  ions  as  calculated 
by  a  simple  rearrangement  of  Equation  4.  This  ratio  in  the  various  sol- 
vents decreases  with  increase  in  dilution  of  the  pairs  constituting  the  cells, 
and  this  in  spite  of  an  increase  in  the  normal  ionization  of  the  electrolyte. 
A  similar  decrease  was  observed  by  Noyes  and  Ellis1  for  concentrated  solu- 
tions of  hydrochloric  acid.  For  concentrations  ranging  from  4.484  N  to 
0.5  N  they  found  a  decrease  in  the  activity  ratios  with  dilution,  followed 
by  an  increase  in  the  value  of  the  ratio  with  further  dilution  between  0.5 
N  and  0.00338  N. 

TABLE  VI. — ACTIVITY  RATIOS  OF  THE  IONS. 

Concentration  ratios  (10  :  1). 


Solvent. 

Water 

A. 
1.0-0.1. 

92  ^4 

B. 
0.5-0.05. 

8  285 

c. 

0.1-0.01. 
7617 

D. 
0.05-0.005, 

6    O42 

Methyl  ale  
Ethyl  ale  . 

6.215 
s  626 

4-725 

4017 

4-03I 

-i    T.QC 

w-Propyl  ale 

464.0 

3147 

3  062 

w-Butyl  ale  

4  *o8i 

4   40^ 

~\   086 

2    6^1 

Isoamvl  ale.  . 

3.QI3 

1.^75 

2.8^ 

2.^00 

The  relative  effect  due  to  an  increase  in  the  concentration  of  the  lithium 
chloride  by  two-,  five-,  ten-,  and  twentyfold  upon  the  activity  ratios  is 
shown  in  Table  VII,  where  A/B,  C/D,  etc.,  represent  the  ratio  of  the  activity 
of  the  ions  for  one  pair  of  concentrations  (A)  to  that  of  another  pair  (B),  etc. 
These  values  point  distinctly  to  an  increase  in  ionic  activity  with  increase 
in  concentration  of  the  dissolved  salt.  As  might  be  expected,  this  increase 
is  least  in  the  aqueous  solutions  and  it  increases  as  we  pass  upward  in  the 
series  of  the  alcohols. 

TABLE  VII. — COMPARATIVE  EFFECT  OF  CONCENTRATION  UPON  THE  ACTIVITY  RATIOS. 

Solvent. 

Water.. 

Methyl  ale 

Ethyl  ale. 

Propyl  ale 

Butyl  ale 

Isoamyl  ale 

1  THIS  JOURNAL,  39,  2543  (1917). 


A/B            C/D            B/C 
2-fold.         2-fold.          5-fold. 

A  1C            B/D 
10-fold.        10-fold. 

A/D 
20-fold. 

I-I7 

.09 

.09 

I  .21 

-19 

1-33 

lc  

17 

•31 

54 

•  *  / 
.22 

•  o  A 

-JQ 

•  OT~ 
7O 

c 

O7 

•  ov 

47 

.  /  *~* 

C  J 

I.  II 

•  **O 

.16 

•T-/ 
4S 

I.6l 

•  O  A 
60 

i  .88 

ilc.  .                           i  .  1  6 

.22 

•  T"  O 

.  JO 

I'.Afl 

.  \-/y 

.46 

i  .60 

i6 

In  the  earlier  part  of  the  discussion  three  factors  were  discussed,  one  or 
more  of  which  may  affect  the  activity  ratios  of  the  ions.  These  were  the 
effect  of  hydration,  the  effect  of  polymerization  of  the  solute  molecules  fol- 
lowed by  complex  ionization  and  the  effect  of  change  in  the  dielectric 
constant  of  the  solution  with  change  in  concentration. 

Lithium  chloride  and  more  especially  the  lithium  ion  exhibit  a  marked 
tendency  to  form  complex  hydrates  in  aqueous  solution.  Similar  solvates 
are  to  be  expected  in  the  various  alcoholic  solutions. 

For  a  salt  which  does  not  form  solvates  the  degree  of  ionization,  and, 
hence,  both  the  concentration  and  the  activity  of  the  ions,  normally  decrease 
with  increase  in  concentration.  For  salts  which  tend  to  form  solvates,  the 
total  amount  of  solvent  removed  from  the  solution  by  solvation  increases 
with  increase  in  the  concentration  of  the  salt,  thus  increasing  somewhat 
abnormally  both  the  concentration  and  the  activity  of  the  ions  for  any  given 
concentration  . 

Considering  any  pair  of  concentrations  constituting  the  cell  and  using 
a  normality  ratio  (10  :  i),  this  abnormal  effect  will  be  greater  in  the  more 
concentrated  solution  and  the  abnormality  will  become  relatively  greater 
as  the  concentration  of  each  of  the  solutions  is  proportionally  increased. 
For  this  reason  the  ratio  of  the  activities  of  the  ions  should  tend  to  increase 
with  increase  in  concentration.  At  sufficiently  high  dilutions,  on  the 
other  hand,  the  effect  due  to  solvation  becomes  a  minimum  and  the  amount 
of  combined  solvent  is  negligibly  small  compared  with  that  actually  present 
as  solvent.  Hence,  increase  in  dilution  should  be  accompanied  by  a  normal 
increase  in  the  concentration  and  activity  ratios. 

The  force  tending  to  bring  about  the  recombination  of  two  oppositely 
charged  ions  is  given  by  the  well-known  relation  of  Coulomb/  =  e^/K.d2. 
Or,  if  we  assume  the  Nernst-Thomson  rule,  for  a  given  concentration  of  the 
electrolyte  in  the  various  solvents,  the  degree  of  ionization,  likewise  the 
concentration  and  the  activity  of  the  ions,  increases  with  the  dielectric 
constant  of  the  solvent. 

Table  VIII  shows  the  relation  between  the  activity  ratios  of  the  ions  and 
the  dielectric  constant  of  the  solvent. 


.  —  RELATION  BETWEEN  ION  ACTIVITY  RATIOS  AND  DIELECTRIC  CONSTANTS.1 

CaHuOH. 
16.0 
4.0 
0.98 
0.84 
0.71 
0.58 

where  Kd  represents  the  dielectric  constant  of  the  solvent  and  A,  B,  C,  D 
represent  the  activity  ratios  of  the  ions  taken  from  Table  VI. 
1  Nernst,  "Theoretical  Chemistry,"  Trans.,  7th  Ger.  ed.,  p.  345. 


H20. 

80  o 

CHsOH. 
1,2    2 

C2HsOH. 
26    I 

CjHyOH. 
22    O 

19  o 

^j_  8-94 
V-Kd  1.03 
\  Kd  o  92 

5.67 
O    QI 

5-108 
O    QI 

4.69 
0.98 

4-356 
1.14 

0.97 

v'jKd-                       o  85 

o  8^ 

O    7O 

o  67 

O   71 

\Kd...                         O.72 

O.7I 

O.6^ 

o.6s 

0.61 

17 

These  data  point  unmistakably  to  the  existence  of  a  definite  relation  be- 
tween the  activity  ratios  and  the  dielectric  constants.  Considering  first 
the  dilute  solutions,  it  will  be  observed  that  the  ratio  (£"/£') /V^,  does  in- 
crease with  increase  in  the  dielectric  constant.  For  all  solvents  this  ratio 
increases  with  increase  in  the  concentration  of  the  solutions  constituting 
the  cells  and  at  sufficiently  high  concentrations  approaches  equal  values 
for  all  solvents.  The  striking  equality  of  this  ratio  for  different  concentra- 
tion ratios  in  water  and  methyl  alcohol  is  significant. 

Hence,  for  sufficiently  high  equivalent  concentrations  of  lithium  chloride 
in  water  and  the  paraffin  alcohols,  we  may  state  that  the  activity  ratios  of 
the  ions  in  equivalent  concentration  ratios  (10  :  i)  are  approximately  pro- 
portional to  the  square  roots  of  the  dielectric  constants  of  the  solvents,  or 
perhaps  better  still,  of  the  solutions.  Unfortunately,  we  do  not  know  the 
dielectric  constants  of  any  of  these  solutions. 

Since  the  activity  ratios  of  the  ions  increase  both  with  the  concentra- 
tion and  with  the  dielectric  constant,  the  activity  of  the  ions  in  a  given  solu- 
tion will  be  greater,  the  greater  both  the  concentration  and  the  dielectric 
constant.  There  is,  therefore,  an  apparently  direct  relation  between  the 
activity  of  the  ions  and  their  electro-affinities.1 

Smale2  has  studied  the  effect  of  various  electrolytes  upon  the  dielectric 
constant  of  water.  Two  of  these  have  an  important  bearing  upon  this 
work.  The  results  are  expressed  in  terms  of  the  dielectric  constant  of 
water  taken  as  unity. 

Salt...     Cone.          o.ooi         0.002         0.005         0.008        o.oio        0.030 
KC1...     K=  1.013        1.018        1.034        1.070        1.113        1.160 

HC1...     K=  0.990        1.033         1.064        1.090        1.126 

Neither  potassium  chloride  nor  its  ions  exhibit  any  appreciable  hydrating 
power.  Hydrogen  chloride,  on  the  other  hand,  shows  a  marked  affinity 
for  water.  Both  increase  the  dielectric  constant  of  water.  The  decrease 
in  the  activity  ratios  of  the  ions  of  potassium  chloride  (Maclnnes  and  Par- 
ker) and  the  increase  in  the  activity  ratios  of  the  ions  of  hydrochloric  acid 
(Noyes  and  Ellis),  both  with  increase  in  concentration,  may  in  all  prob- 
ability be  due  entirely  to  differences  in  hydrating  power.  The  activity 
ratios  of  the  ions  should  increase  with  the  concentration  most  rapidly  for 
salts  which  have  high  dielectric  constants  and  great  hydrating  power. 

It  is  obvious,  therefore,  that  changes  in  the  free  energy  of  dilution,  of 
electromotive  force  and  other  factors  depending  upon  the  activity  ratios 
must  be  intimately  related  to  changes  in  the  dielectric  constants  of  the 
solutions. 

While  the  effects  due  to  solvation  and  changes  in  dielectric  constant  may 

1  By  electro-affinity  is  meant  the  power  of  a  cation  to  repel,  or  the  power  of  an 
anion  to  attract  an  electron. 

2  Ann.  Phys.  Chem.,  [ii]  61,  625  (1897). 


i8 

explain  in  part  the  increase  in  the  activity  ratios  for  solutions  in  water  and 
the  lower  alcohols,  they  alone  will  not  suffice  for  the  interpretation  of  the 
effects  observed  in  the  higher  alcohols.  In  the  latter  we  undoubtedly  have 
present  not  only  the  simple  molecules  and  simple  ions,  but  also  polymerized 
molecules  and  complex  ions,  all  in  equilibrium  as  represented  in  (50)  or 
(56).  A  continuous  increase  in  the  concentration  of  the  salt  results  in  a 
constantly  increasing  concentration  of  one  of  the  complex  ions,  e.  g.t 
LiCl2~.S5,  or  I42C1+.S6.  There  is  also  a  corresponding  increase  in  the 
number  of  the  oppositely  charged  simple  ion.  The  activity  ratio  of  the 
negative  or  positive  ions  now  involves  the  activities  of  both  the  simple  and 
complex  ions  of  each. 

For  any  concentration  ratio  (10  :  i)  the  effect  produced  by  polymeriza- 
tion will  always  be  greater  in  the  more  concentrated  solution  and  it  will 
increase  relatively  the  more  rapidly  the  greater  the  proportional  increase 
in  the  concentration  of  the  two  solutions  constituting  the  cell.  The  ab- 
normal increase  in  the  concentration  of  the  complex  ions,  coupled  with  an 
increase  in  the  dielectric  constant  and  with  an  increase  in  the  total  amount  of 
combined  solvent,  with  increase  in  concentration  of  the  salt,  should  and  do 
increase  the  activity  ratios  of  the  ions. 

TABLE  IX. — ACTIVITY  RATIOS  OF  THE  MOLECULES. 

Concentration  ratios  (10  :  1). 


Solvent. 
Water  

1.0-0.1. 
85.63 

0.5-0.05. 
68.64 

0.1-0.01. 
58.02 

0.05-0.005. 
48.19 

Methyl  ale  

38.63 

22.33 

16.25 

Ethyl  ale    

31  .65 

I6.3O 

13.75 

w-Propyl  ale 

21    5^ 

9QO 

9.37 

n-Butyl  ale  

24..  81 

20.  19 

9.52 

7.03 

Isoamvl  ale.  .  , 

1.5.  3  1 

1  1  .  39 

8.03 

5-30 

The  effect  of  solvent  and  concentration  upon  the  ratio  of  the  molecules 
is  shown  in  Table  IX.  From  the  relation  presented  in  Equation  4  these 
ratios  are  equal  to  the  square  of  the  corresponding  activity  ratios  of  the  ions. 
Factors  influencing  the  latter  will  also  influence  these  accordingly.  These 
data  require  no  further  discussion. 

Summary. 

The  electromotive  forces  of  concentration  cells  have  been  determined 
for  solutions  of  lithium  chloride  in  water,  methyl,  ethyl,  w-propyl,  w-butyl 
and  isoamyl  alcohols.  All  of  the  cells  measured  contained  solutions  having 
a  normality  ratio  of  10  :  i. 

The  electromotive  force  with  transference  increases  with  increasing  dilu- 
tion for  all  of  the  solvents  used,  except  butyl  and  isoamyl  alcohols  in  which  it 
decreases.  The  electromotive  force  without  transference  decreases  with 
increases  in  dilution  in  all  of  the  solvents. 


19 

The  transport  number  of  the  lithium  ion  has  been  calculated  and  it  has 
been  found  to  increase  in  each  solvent  as  the  concentration  of  the  salt  is 
diminished. 

The  free  energy  of  dilution  and  the  activity  ratios  of  both  the  ions 
and  the  molecules,  on  the  other  hand,  decrease  with  the  dilution. 

An  attempt  has  been  made  to  explain  the  decrease  in  the  activity  ratios, 
and  hence  the  free  energy,  on  the  assumption  of  eff ects  due  to  hydration  and 
the  change  in  dielectric  constant.  For  the  higher  alcohols,  it  has  been 
found  necessary  to  assume  a  polymerization  and  subsequent  complex  ioni- 
zation  of  the  salt  molecules. 

For  a  given  pair  of  concentrations  the  activity  ratios  of  the  ions  increase 
with  increase  in  the  dielectric  constant  of  the  solvent.  The  ratio 
(£" ' /%')/^Kd  increases  with  increase  in  concentration  and  at  the  same  suffi- 
ciently high  equivalent  concentrations  attains  an  approximately  equal 
value  for  all  solvents.  For  these  concentrations  the  ration  (£"/£')  is  very 
nearly  directly  proportional  to  the  square  root  of  the  dielectric  constant 
of  the  solvent.  For  similar  cells  in  water  and  methyl  alcohol  these  ratios 
are  practically  identical. 

The  activity  of  the  ions  has  been  found  to  increase  both  with  the 
concentration  and  with  the  dielectric  constant  of  the  solvent,  or  of 
the  solution. 


BIOGRAPHY. 

Franklin  Spencer  Mortimer  was  born  in  Valton,  Wisconsin,  May  20, 
1891.  He  received  his  early  education  in  the  public  schools  of  that  vil- 
lage. He  entered  Penn  Academy,  Oskaloosa,  Iowa,  in  1908.  In  1914 
he  was  graduated  from  Penn  College  with  the  degree  of  Bachelor  of  Science. 

In  the  Fall  of  1914  he  entered  the  State  University  of  Iowa,  where  he 
pursued  graduate  work  in  Chemistry.  His  major  work  was  in  the  De- 
partment of  Physical  Chemistry,  while  Industrial  Chemistry  and  Geology 
constituted  his  minor  subjects. 

He  was  appointed  to  a  Fellowship  for  the  year  1916-17  by  the  Graduate 
College.  In  the  Spring  of  1916  he  had  the  honor  of  being  elected  a  mem- 
ber of  the  Society  of  the  Sigma  Xi. 

He  is  now  Professor  of  Chemistry  in  Penn  College,  Oskaloosa,  Iowa. 


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ArR    7    1934 

APR    8     1934 

i    to^B.  =-' 

\8lu<v5,p" 

^ 

- 

LD  21-100m-7,'33 

Binder 
Gaylord  Bros. 

Makers 
Syracuse,  N.  y 

MT.  JAN  2J,  1905    * 


545. '121 


M* 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 


